Atomic Structure and the Periodic Table
This domain tests your understanding of what atoms are made of, how electrons are arranged, and how position in the Periodic Table follows from that arrangement. You must be able to use proton number, mass number, isotopes and electronic configuration confidently.
Subatomic particles and notation
- A proton has relative mass 1 and charge +1; a neutron has relative mass 1 and charge 0; an electron has negligible relative mass and charge –1.
- Proton number (atomic number) equals the number of protons and, in a neutral atom, the number of electrons. Mass number equals protons plus neutrons.
- Isotopes are atoms of the same element with the same proton number but different numbers of neutrons. They have identical chemical properties but slightly different physical properties and masses.
Electron arrangement and the Periodic Table
Electrons occupy shells in the order 2, 8, 8. The number of outer-shell electrons equals the group number for main-group elements (Group 0 has a full outer shell). The number of occupied shells equals the period number. Metals lie to the left and centre; non-metals to the right.
Worked example
An atom has 17 protons, 18 neutrons and 17 electrons. Deduce its mass number, electronic configuration, group and period, and the charge on its ion. Mass number = 17 + 18 = 35. Configuration = 2,8,7. Seven outer electrons place it in Group VII (the halogens); three occupied shells place it in Period 3. It needs one more electron to complete its outer shell, so it forms a 1– ion.
Traps that cost marks
- Confusing mass number with relative atomic mass; relative atomic mass is a weighted average of isotope masses.
- Writing electron configurations such as 2.8.7 instead of 2,8,7.
- Saying isotopes have different chemical properties; they do not.
- Assuming all ions have the same charge as their group number; metals lose electrons, non-metals gain them.